Lewis Structures and Formal Charge: A Method You Can Check

Counting valence electrons for molecules and polyatomic ions

Every Lewis structure starts with one number: the total count of valence electrons available to place. For main-group elements, the valence electron count equals the last digit of the group number in the 1–18 numbering, so carbon contributes 4, nitrogen 5, oxygen 6, and each halogen 7. Hydrogen contributes 1, and helium, with 2, is the one exception to the group-number rule.

An ion's charge changes the total. A negative charge means extra electrons, so you add one electron per unit of negative charge. A positive charge means missing electrons, so you subtract. The ammonium ion, NH₄⁺, has 5 + 4(1) − 1 = 8 valence electrons, while the carbonate ion, CO₃²⁻, has 4 + 3(6) + 2 = 24.

Write this total down before you draw anything; it is the number you will check the finished structure against.

Building the skeleton and placing electron pairs

The skeleton is the pattern of which atom connects to which. Hydrogen and fluorine almost always form only one bond, so they are terminal atoms. The central atom is usually the least electronegative atom other than hydrogen, which is often the atom that appears once in the formula. In oxyacids, hydrogen atoms are normally bonded to oxygen, not to the central atom.

Once the skeleton is set, the remaining steps distribute electrons so that each atom reaches a stable count. Period 2 atoms (C, N, O, F) can hold at most eight electrons around them because they have only four valence orbitals, which is why the octet rule is a firm limit for them.

• Connect each terminal atom to the central atom with a single bond (2 electrons each).
• Complete the octets of the terminal atoms with lone pairs (hydrogen needs only 2).
• Place any electrons that remain on the central atom as lone pairs.
• If the central atom still has fewer than 8 electrons, convert a lone pair on a neighboring atom into a shared pair to make a double or triple bond.
• Check: electrons drawn equals the total counted, and each period 2 atom has no more than 8.

Formal charge as a test between candidate structures

Formal charge asks how many electrons an atom owns in the drawing compared with the neutral free atom, if every bonding pair were split evenly. The formula is FC = V − N − B/2, where V is the atom's valence electron count, N is the number of nonbonding (lone-pair) electrons on it, and B is the number of electrons in its bonds. The formal charges in a structure always add up to the overall charge of the species, which gives you a built-in arithmetic check.

When several structures obey the electron count, prefer the one in which formal charges are as close to zero as possible. If some formal charge must remain, the better structure places negative formal charge on the more electronegative atom and avoids putting charges of the same sign on bonded neighbors.

Formal charge is bookkeeping, not a measured charge. Because it ignores electronegativity, an atom with zero formal charge can still carry a substantial partial charge in the real molecule.

Worked example: phosgene and the cyanate ion

Phosgene, COCl₂. Count: C contributes 4, O contributes 6, and each Cl contributes 7, so the total is 4 + 6 + 2(7) = 24 electrons. Carbon is the central atom. Three single bonds from C to O and to each Cl use 6 electrons. Completing the octets of O and both Cl atoms uses 3 × 6 = 18 more, for 24 in all, so no electrons remain for carbon. Carbon has only 6 electrons around it, so convert one lone pair into a second bond. Moving an O lone pair gives C=O; moving a Cl lone pair gives C=Cl.

Formal charge decides between them. With C=O, oxygen has 4 nonbonding electrons and 4 bonding electrons: FC = 6 − 4 − 4/2 = 0. Each Cl has 6 nonbonding and 2 bonding electrons: FC = 7 − 6 − 1 = 0. Carbon has 8 bonding electrons: FC = 4 − 0 − 4 = 0. With C=Cl instead, that Cl becomes 7 − 4 − 2 = +1 and the singly bonded O becomes 6 − 6 − 1 = −1. The all-zero structure with the C=O double bond is preferred.

Cyanate ion, OCN⁻. Count: 6 + 4 + 5 + 1 = 16 electrons, with carbon central. Three arrangements give carbon an octet: O=C=N, O≡C–N, and O–C≡N, each with lone pairs completing the octets on O and N. For O=C=N, O = 6 − 4 − 2 = 0 and N = 5 − 4 − 2 = −1. For O≡C–N, O = 6 − 2 − 3 = +1 and N = 5 − 6 − 1 = −2. For O–C≡N, O = 6 − 6 − 1 = −1 and N = 5 − 2 − 3 = 0. Carbon is 0 in all three, and each set sums to −1 as it must.

The O≡C–N form has the largest charges and puts positive charge on oxygen, the most electronegative atom, so it contributes least. The O–C≡N form places the negative charge on oxygen and is the most important contributor, with O=C=N close behind. The real ion is a blend of these structures, as the next section explains.

Resonance when one drawing cannot describe the bonds

Some species have two or more valid Lewis structures that differ only in where electrons are placed, with every atom in the same position. These are resonance structures, written with a double-headed arrow (↔) between them. The real molecule is a single structure, the resonance hybrid, whose bonds are an average of the contributors. It does not switch back and forth between them.

The formate ion, HCOO⁻, is a clean example. Count: 1 + 4 + 2(6) + 1 = 18 electrons. With H and both O atoms bonded to C and the oxygen octets completed, 18 electrons are used and carbon has only 6, so one oxygen lone pair becomes a C=O bond. The double-bonded O has FC = 6 − 4 − 2 = 0 and the single-bonded O has FC = 6 − 6 − 1 = −1. Either oxygen could hold the double bond, so there are two equivalent structures.

Because the two contributors are equivalent, the hybrid has two identical carbon–oxygen bonds, each intermediate between a single and a double bond, and each oxygen carries half of the negative charge on average. Measured bond lengths in carboxylate ions agree with this picture. When contributors are not equivalent, as in cyanate, the hybrid resembles the lower-formal-charge structures most.

When the octet rule bends: odd, incomplete and expanded counts

Three kinds of species do not fit the octet pattern. Odd-electron molecules, such as nitrogen monoxide, NO, with 5 + 6 = 11 valence electrons, cannot give every atom an octet; one atom must carry an unpaired electron. Species of this kind are called radicals and are typically reactive.

Some molecules have an incomplete octet on the central atom. In boron trifluoride, BF₃ (24 electrons), the structure with three B–F single bonds leaves boron with 6 electrons but gives every atom a formal charge of zero. Forming a B=F double bond would complete boron's octet only by giving fluorine a formal charge of +1 and boron −1, so the single-bonded structure is the usual description.

Atoms in period 3 and below can be drawn with more than eight electrons around them. In PCl₅ (40 electrons), five bonds give phosphorus 10 electrons. In the triiodide ion, I₃⁻ (3 × 7 + 1 = 22 electrons), the central iodine holds two bonds and three lone pairs, and its formal charge is 7 − 6 − 2 = −1. Place leftover electrons on the central atom only after every terminal atom has its octet.

Common Lewis structure errors and how to catch them

Most mistakes are counting mistakes, so the most reliable habit is to finish every structure with two checks: count every electron in the drawing (2 per bond, 2 per lone pair) against your written total, and add the formal charges to confirm they equal the overall charge.

If your total differs from a classmate's or the key by twice the ion's charge, the charge was probably applied with the wrong sign. If the drawing holds a different number of electrons than your total, look for bonds counted as one electron instead of two, or for leftover electrons that never got placed. The specific errors below account for most wrong structures.

• Adding electrons for a cation or subtracting them for an anion. Negative charge means more electrons.
• Counting a bond as one electron. Each line represents a shared pair, so it is 2 electrons.
• Giving hydrogen two bonds or making it the central atom. Hydrogen holds only 2 electrons.
• Placing more than 8 electrons on C, N, O or F. Expanded counts are possible only for period 3 and lower.
• Treating formal charge as the real charge on an atom, or confusing it with oxidation number, which assigns all bonding electrons to the more electronegative atom.
• Drawing resonance structures by moving atoms. Moving an atom creates a different compound, not a resonance contributor.